Reaction Types: Synthesis, Decomposition, Replacement, and Combustion

Why classify reactions?

When you look at a chemical equation, it can feel like a new language. Reaction “types” are a way to recognize common patterns so you can predict products, check whether an equation makes sense, and connect the equation to what you would observe in a lab or in everyday life (gas bubbles, heat, a new solid forming, and so on). In this chapter you will learn four very common patterns: synthesis, decomposition, replacement (single and double), and combustion.

Important note: a “type” is a pattern, not a law. Some reactions fit more than one category, and some real-world reactions are mixtures of steps. Still, these four patterns cover a large fraction of beginner-level chemistry problems.

1) Synthesis (Combination) Reactions

Core idea and pattern

A synthesis reaction forms one main product from two (or more) simpler reactants. The pattern is:

A + B → AB

Think “pieces combine to make a single compound.” Often, synthesis reactions release energy (but not always), and they commonly occur when elements react to form compounds or when simple compounds combine into a more complex compound.

How to recognize a synthesis reaction

• There are multiple reactants and one main product.
• The product is typically a compound (not always, but commonly).
• The reactants may be elements, or small compounds like oxides.

Common synthesis sub-patterns you will see

1) Element + element → ionic compound

Continue in our app.

You can listen to the audiobook with the screen off, receive a free certificate for this course, and also have access to 5,000 other free online courses.

Or continue reading below...

Download the app

2Na + Cl2 → 2NaCl

Here, a metal and a nonmetal combine. In many problems, the product formula is determined by the typical charges of the ions formed. (You do not need to re-derive those charge rules here; just apply them.)

2) Metal oxide + water → metal hydroxide (a base)

CaO + H2O → Ca(OH)2

This pattern is important in materials and construction chemistry. Calcium oxide (quicklime) reacts with water to form calcium hydroxide (slaked lime), releasing heat.

3) Nonmetal oxide + water → oxyacid (an acid)

SO3 + H2O → H2SO4

Many acid rain chemistry discussions involve this type of synthesis: sulfur oxides and nitrogen oxides reacting with water to form acids.

4) Metal + oxygen → metal oxide

2Mg + O2 → 2MgO

Burning magnesium ribbon is a classic example: bright white light and a white powder (magnesium oxide) form.

Practical step-by-step: predicting products for synthesis

When you are given reactants and asked to predict the product, use this checklist.

• Step 1: Confirm it is “many reactants → one product.” If yes, synthesis is likely.
• Step 2: Identify what kind of reactants you have. Element + element? Oxide + water? Something else?
• Step 3: Write the likely product formula. For ionic products, use ion charges to get the correct ratio. For oxide + water patterns, use the standard forms: metal hydroxide or oxyacid.
• Step 4: Sanity-check the product. Does it contain all the elements present in the reactants? In synthesis, it should.

Example (step-by-step): Predict the product of aluminum and oxygen.

• Step 1: Two reactants, one product expected → synthesis.
• Step 2: Metal + oxygen → metal oxide.
• Step 3: Aluminum forms Al3+ and oxygen forms O2−, so the oxide is Al2O3.
• Step 4: Product contains Al and O only → good.

Al + O2 → Al2O3

(Balancing is a separate skill you already practiced; here the focus is identifying the type and predicting the product.)

2) Decomposition Reactions

Core idea and pattern

A decomposition reaction breaks one compound into two (or more) simpler substances. The pattern is:

AB → A + B

Think “one thing splits into parts.” Decomposition often requires an input of energy such as heat, electricity, or light, because breaking bonds typically takes energy.

How to recognize a decomposition reaction

• There is one reactant and multiple products.
• The reactant is usually a compound.
• The products are simpler: elements and/or smaller compounds.

Common decomposition sub-patterns

1) Binary compound → elements

2HgO → 2Hg + O2

Mercury(II) oxide decomposes into mercury and oxygen when heated.

2) Metal carbonate → metal oxide + carbon dioxide

CaCO3 → CaO + CO2

This is central to cement and lime production. Heating limestone (calcium carbonate) produces calcium oxide and carbon dioxide.

3) Metal chlorate → metal chloride + oxygen

2KClO3 → 2KCl + 3O2

Potassium chlorate decomposes to release oxygen gas (often with a catalyst such as MnO2). The oxygen supports combustion strongly.

4) Hydrogen peroxide → water + oxygen

2H2O2 → 2H2O + O2

This can happen slowly on its own, or quickly with catalysts (like the enzyme catalase). The bubbling you see is oxygen gas.

Practical step-by-step: predicting products for decomposition

Because decomposition starts with one compound, the key is recognizing what “family” the compound belongs to.

• Step 1: Confirm the pattern “one reactant → multiple products.”
• Step 2: Identify the compound type. Is it a carbonate (contains CO3)? A chlorate (contains ClO3)? A peroxide (contains O2 as a unit, like H2O2)?
• Step 3: Apply the matching decomposition rule. Carbonate → oxide + CO2; chlorate → chloride + O2; peroxide → oxide (or water) + O2 depending on the peroxide.
• Step 4: Check that atoms are conserved in your unbalanced equation. You should be able to account for every element from the reactant in the products.

Example (step-by-step): Predict the products when sodium carbonate decomposes.

• Step 1: One reactant → decomposition.
• Step 2: Contains CO3 → carbonate.
• Step 3: Carbonate → metal oxide + CO2.
• Step 4: Sodium carbonate (Na2CO3) becomes sodium oxide (Na2O) and carbon dioxide (CO2).

Na2CO3 → Na2O + CO2

3) Replacement Reactions

Replacement reactions involve swapping partners. They are extremely useful because they connect the equation to a simple “who pairs with whom” idea. There are two main kinds: single replacement and double replacement.

3A) Single Replacement (Single Displacement)

Core idea and pattern

In a single replacement reaction, one element replaces another element in a compound. The pattern is:

A + BC → AC + B

or, if a nonmetal like a halogen replaces another halogen:

X2 + BY → BX + Y2

These reactions are driven by relative reactivity: some elements more readily form ions (or gain electrons) than others. In practice, you often use an activity series (for metals) or halogen reactivity trends to decide whether the reaction happens.

How to recognize single replacement

• One reactant is an element (a pure metal like Zn, or a diatomic halogen like Cl2).
• The other reactant is a compound.
• Products show the element inserted into the compound and a different element released.

Metal replacement: typical observations

When a metal replaces another metal in a solution, you may see the “new” metal appear as a coating or solid, and the solution color may change if ions change. For example, copper(II) solutions are often blue; if copper metal forms, you may see reddish-brown copper deposit.

Practical step-by-step: deciding if a single replacement occurs

• Step 1: Identify the free element. Is it a metal (like Mg) or a halogen (like Br2)?
• Step 2: Identify what it could replace. In A + BC, A can only replace B (a similar type: metal replaces metal; halogen replaces halogen).
• Step 3: Use a reactivity rule to decide if it happens. For metals, a more reactive metal can replace a less reactive metal from its compound. For halogens, a more reactive halogen can replace a less reactive halogen from a halide.
• Step 4: If it happens, write products by swapping and keeping charges consistent. The new compound AC must have the correct formula based on ion charges.
• Step 5: If it does not happen, write “no reaction.” Many textbook questions test this decision.

Example (step-by-step): Predict whether zinc metal reacts with copper(II) sulfate solution.

• Step 1: Free element is Zn (a metal).
• Step 2: It could replace Cu in CuSO4.
• Step 3: Zinc is more reactive than copper (zinc more readily forms Zn2+), so replacement occurs.
• Step 4: Products: ZnSO4 and Cu.

Zn + CuSO4 → ZnSO4 + Cu

Example where no reaction occurs: Copper metal placed in magnesium sulfate solution.

• Cu is less reactive than Mg, so it cannot replace Mg2+.

Cu + MgSO4 → no reaction

Halogen replacement: a common pattern

Halogens (F2, Cl2, Br2, I2) can replace each other in salts. Reactivity generally decreases down the group: F2 is most reactive, then Cl2, then Br2, then I2. So chlorine can replace bromide, but bromine cannot replace chloride.

Cl2 + 2KBr → 2KCl + Br2

Br2 + 2KCl → no reaction

3B) Double Replacement (Double Displacement / Metathesis)

Core idea and pattern

In a double replacement reaction, two ionic compounds exchange ions. The pattern is:

AB + CD → AD + CB

This is like “partner swapping.” These reactions commonly occur in aqueous solution (dissolved in water), where ions can move freely and recombine into new compounds.

When does a double replacement reaction actually happen?

If you simply swap ions on paper, you can always write “products,” but in real chemistry the reaction proceeds only if there is a driving force that removes ions from solution. The most common driving forces at this level are:

• Formation of a precipitate: an insoluble solid forms and separates from the solution.
• Formation of a gas: bubbles form and leave the solution.
• Formation of water: often in acid-base neutralization, water forms and is a very stable product.

You do not need a full solubility table memorized to start recognizing typical precipitates, but you should know that many nitrates are soluble, many alkali metal salts are soluble, and many carbonates and hydroxides are insoluble except with alkali metals. In practice problems, you are often told a precipitate forms, or you are given compounds that commonly form insoluble products (like AgCl).

Practical step-by-step: predicting products for double replacement

• Step 1: Confirm you have two compounds as reactants. Often they are ionic and may be labeled (aq).
• Step 2: Swap the positive ions (cations) between the negative ions (anions). Keep the polyatomic ions intact as units.
• Step 3: Write correct formulas for the new compounds. Ensure charges balance in each product.
• Step 4: Decide if a driving force exists. Look for precipitate, gas, or water formation. If none, it may be “no reaction” in aqueous solution.

Example (precipitation): Mix aqueous silver nitrate and aqueous sodium chloride.

• Step 1: Two ionic compounds: AgNO3 and NaCl.
• Step 2: Swap partners: Ag+ pairs with Cl−; Na+ pairs with NO3−.
• Step 3: Products: AgCl and NaNO3.
• Step 4: AgCl is insoluble (forms a solid precipitate), so reaction proceeds.

AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Example (gas formation): Acid + carbonate often produces carbon dioxide gas.

2HCl(aq) + Na2CO3(aq) → 2NaCl(aq) + H2O(l) + CO2(g)

Here the “swap” idea gets you to carbonic acid (H2CO3) as an intermediate, which quickly decomposes into H2O and CO2. In practice, you will often write the final products directly as water and carbon dioxide because that is what is observed (bubbling).

Example (water formation / neutralization):

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Even if the salt remains dissolved, the formation of water drives the reaction.

4) Combustion Reactions

Core idea and pattern

Combustion is a reaction with oxygen that releases energy as heat (and often light). In beginner chemistry, “combustion reaction” usually means a substance reacts with O2 to form oxides, and for hydrocarbons (compounds containing only carbon and hydrogen) the products are carbon dioxide and water.

Hydrocarbon combustion pattern:

CxHy + O2 → CO2 + H2O

If the fuel contains oxygen as well (like ethanol, C2H5OH), the products are still typically CO2 and H2O in complete combustion:

C2H5OH + O2 → CO2 + H2O

Complete vs incomplete combustion (what changes and what you might observe)

Complete combustion happens when there is plenty of oxygen. Carbon ends up as CO2, and hydrogen ends up as H2O. Flames are often blue and cleaner, and less soot forms.

Incomplete combustion happens when oxygen is limited. Carbon may form carbon monoxide (CO) and/or solid carbon (soot). Flames may appear yellow/orange due to glowing soot particles, and you may see black deposits.

Examples of incomplete combustion product patterns include:

CxHy + O2 → CO + H2O

CxHy + O2 → C(s) + H2O

In real situations, incomplete combustion often produces a mixture of CO2, CO, H2O, and soot.

Practical step-by-step: writing products for combustion

Step 1: Identify whether the reactant is a hydrocarbon or an oxygen-containing organic compound. If it is mostly C and H (and maybe O), combustion questions usually expect CO2 and H2O as products for complete combustion.

Step 2: Write CO2 and H2O as products. Do not guess unusual products unless the problem states “incomplete combustion” or indicates limited oxygen.

Step 3: Add O2 as a reactant if it is not already present. Combustion requires oxygen gas.

Step 4: (If needed) Determine whether it is complete or incomplete. If the problem mentions soot, yellow flame, or CO, treat it as incomplete.

Example (complete combustion): Propane burns in oxygen.

C3H8 + O2 → CO2 + H2O

Example (incomplete combustion): Methane burns with limited oxygen, producing carbon monoxide.

CH4 + O2 → CO + H2O

Combustion beyond hydrocarbons: metals and other substances

Combustion can also describe rapid oxidation of metals. For example, magnesium reacts vigorously with oxygen to form magnesium oxide:

2Mg + O2 → 2MgO

This is both a synthesis reaction (two reactants form one compound) and a combustion/oxidation reaction (reaction with oxygen releasing energy). Classification depends on what you are focusing on.

Putting it together: a quick identification workflow

When you are given an unfamiliar reaction, use this decision path to classify it quickly.

• 1 reactant? Likely decomposition (AB → A + B).
• 2+ reactants and 1 main product? Likely synthesis (A + B → AB).
• Element + compound? Likely single replacement (A + BC → AC + B), but check whether it actually occurs based on reactivity.
• Compound + compound? Likely double replacement (AB + CD → AD + CB), but check for precipitate/gas/water driving force.
• O2 is a reactant and energy release is implied? Likely combustion. For hydrocarbons, expect CO2 and H2O (complete combustion).

Practice set with guided hints (do on paper)

Classify each reaction type

• 1) 2K + Br2 → 2KBr

  Hint: two reactants, one product.

• 2) 2NaN3 → 2Na + 3N2

  Hint: one compound breaks apart (this is related to airbag chemistry).

• 3) Fe + CuSO4 → FeSO4 + Cu

  Hint: element + compound; ask which metal is more reactive.

• 4) BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + 2NaCl(aq)

  Hint: two ionic compounds swap; look for a solid.

• 5) C4H10 + O2 → CO2 + H2O

  Hint: hydrocarbon + oxygen.

Predict products (type first, then products)

• A) Al + Cl2 → ?

  Hint: synthesis; ionic compound forms.

• B) K2CO3 → ?

  Hint: carbonate decomposition.

• C) Cl2 + NaI → ?

  Hint: halogen replacement; compare halogen reactivity.

• D) HNO3(aq) + KOH(aq) → ?

  Hint: double replacement; water formation.

• E) C2H5OH + O2 → ?

  Hint: complete combustion gives CO2 and H2O.